The Acidic Environment
Describe the solubitliy of carbon dioxide in water in terms of LCP AND describe changes with conc, addition fo acid or base, and gas levels
1. the CO2 dissolving into the water CO2(g) <-> CO29aq) 2. the formation of carbonic acid CO2(g) + H2O(l) ⇌ H+(aq) + HCO3-(aq)ΔH < 0 3. 2 step dissociation of the carbonic acid H2CO3(aq) + H2O(l) ⇌ H3O+(aq) + HCO3-(aq) ΔH < 0 HCO3-(aq) + H2O(l) ⇌ H3O+(aq) + CO32-(aq) ΔH < 0
What is a neutralisation reaction
A neutarlisation reaction is when an acid and base react together to form a salt and water. Exothermic reaciton and irreverisble/completion
Describe the differences between the alkanol and alkanoic acids in terms of structure. Give the general formlua as well
ALKANOL: * contains the OH group i.e. the hydroxyl group * CnH2n+1OH ALKANOIC ACID * contains the carboxylic functional group * CnH2n-1COOH
Describe the pH of metal oxides and non metal oxides and where they are located in the periodic table. Give examples of in water and then neutraliastion
Acidic oxides * Most non metal oxides are acidic oxides * They ionise in water to produce hydronium ions E.g. CO2 + H2O <-> H2CO3 H2CO3 + H2O <-> H3O+ + HCO3- CO2 + 2NaOH --> H2O + Na2CO3 * located to the right side since mainly non metal oxides * increase across the period, decreasing from top to bottom fo the group METALLIC OXIDES * metal oxides are basic thereofre occurs mainly left side E.g. CaO + H2O <-> Ca2+ + 2OH- MgO + H2O <-> Mg2+ + 2OH- CuO + H2SO4 --> CuSO4 + H2O
Define acid and base and how they ionise in water
Acids * The Bronsted Lowry definition of an acid is a substance which produces hydrogen ions (H+) or donates protons * In water, acids ionise and react with water to form hydronium ions (H3O+) or hydrated hydrogen ion as free H+ does not exist in an aqueous solution Bases * The BL definition of a base ia substance which produces hydroxide (OH-) or accepts protons * In water bases ionise and react with water to form hydroxide ions
Gather and process information from secondary sources to explain the use of acids as food additives
Acids are used as food additives to: * lwoer the pH so that the enzyme activity is inhibited --> prevents the growth of microbes which can spoil the food * improve the flavour by adding tartness * enhance nutritaionl value E.g. ascorbic acid C6H8O6 used as a preservative and improves Vit C content Citric acid is used in soft drinks as a preservative and can improve flavour Acetic acid is used as a preservative in picking and is used to improve flavour
Give examples of amphoteric substances and where they are located in the periodic table
Act as an acid and base therefore can neutralise both acids and bases Be, Al, Zn, Sn and Pb. Located between the borderline of metals and non metals e.g. Aluminium oxide Al2O3 + 6HCl <--> AlCl3 + H2O (balance lmao) Al2O3 + 2NaOH <-> 2NaAlO2 + H2O
How to determine the acidic, neutral or basic nature of solutoisn?
Hydrolysis fo slats and the ions formed need to be considred separately (hydrolysis is the reaction fo a salt and water to produce a change in pH) HCl + NH3 —> NH4Cl H+ + Cl- + NH3 —> NH4+ + Cl- * NH4Cl is weakly acidic as it is a weak base + strong acid * Cl- is neutral and NH4+ is a weak acid as NH4+ + H2O ⇌ NH3 + H3O+ * I.e. it has donated protons = proves it is acidic because ions are proton donors NaCH3COO * Slightly basic * NaCH3COO —> Na+ + CH3COO- * This is because CH3CO O- is a strong base as CH3COO- + H2O ⇌ CH3COOH + OH- (remember equilibrium symbol as it is a weak base in water, and also CH3COO- is weak as it is formed from a weak acid CH3COOH) * Na+ is neutral and CH3COO- is slightly basic as it can accept protons * IF IT IS A SPECTATOR ION = NEUTRAL
What is a functional group
It is an atom or a group of atoms which change the chemical properties of a molecule
Gather and process information from secondary sources to trace developments in understanding and describing base/acid reactions
LAVOISER * Proposed that acids contain oxygen since many non metal oxides (CO2 and SO2) produced an acidic solution when reacted with water. This was a component in the compound responsible for the generic properties of acids, the other was called an acidifiable base * However this acidifying principle did not explain why HCl and HF were acidic, or why metal oxides were not acidic DAVY * he proposed that an acid contained replacebale hydrogen i.e. when reacted with a metal, it will be replaced by the metal and thus form a metal salt and H2 * he disproved Lavoisier by experimenting with muriatic acid with various metals and non metals, and he never obtained any oxygen compounds. When reacted with mercury, he found it produced a metal salt and hydrogen gas * he further disproved that oxygen was the acidifying principle by showing that alkali metals and alkali earth metals were oxides e.g. Mg, Ca and Ba * however this did not explain why e.g. CH4 was not acidic Metal + acid --> H2 + metal salt ARRHENIUS * he proposed that acids ionise in water to form H3O+ and bases ionise in water to form OH- ions * he discovered that an acid's degree of ionisation was related to its strength * this also explains the process of neutralisation, where teh OH- ions int he base react with teh H+ ions from the acid to form H2O Acid + base --> salt + water * however this did not explain the role of the solvent in the ionisatino fo the acid i.e. acids cannot be ionised without the presence of solvent * didn't explain amphoteric substances * didn't explain the acidic or basic nture of substances even though they don't release H3O or OH- ions BL * * The BL theory of acids and bases: * An acid is a proton donor * A base is a proton acceptor * A proton being an H+ ion * By redefining acid base reactions as proton-transfer reactions and not just as reactions in water, this theory can explain: Why his was superior * could explain amphotheric substances * why some acids and bases do not contain hydrogen or OH- ions * extended to non aqueous media e.g solids and gases * could allow for acids to include positive and negative ions as well as neutral molecuels * expalint he difference int he relative strengths between acdis and between bases * explain the realtionshpi ebtween the strength of an acid and its conjugate base * explain the role of water in the acid base reaciotn i.e. accepts H+ ions when reacted with acid * expanded the list of abses to include any moleucle or iont aht cna accept a proton * can explain the levelling effect in water -- strong acids and bases allhave the save strength when dissolved in water Example NH3 + HCl --> NH4Cl *According to Arrhenius theory the above reaction is not an acid/base neutralisation as they are not in aqueous solutions, no free hydrogen ions or hdyroxide ions released and no water is formed when neturalised * Also NH3 is not considered an Arrhenius base as it does not contain the OH- ions that must be ionised into the water * According to the BL theory, the HCl molecule has donated a proton and the NH3 molecule has accepted a proton. As a transfer of protons has occurred, the HCl molecule is a BL acid and the NH3 molecule is.a BL base
pH range of salt solutions experiment
pH meter (review the instructions for use from Experiment 7) • Small beakers for each solution. • Spotting plates CHEMICALS Universal indicator Approximately 45 mL of each of the following salts should be tested: 0.1M ammonium chloride 0.1M potassium nitrate 0.1M sodium acetate 0.1M sodium carbonate 0.1M sodium phosphate 0.1M ammonium acetate 0.1M ammonium sulfate 0.1M sodium chloride 0.1M sodium hydrogen carbonate distilled water Risk assesment Method 1. Place a small amount of each salt solution in a separate semi micro test tube or on a spotting plate and test for pH using Universal Indicator. 2 Use a pH meter to determine the pH and record your results in the results table. 3 Be sure to rinse the electrode thoroughly between measurements. TABLE: salt, formular, pH using universal indicator, pH meter reading, acid, basic or neutral
What are the aspects of a closed system
* Closed (no matter or energy can leave the system) * Macroscopic properties stay constant (macroscopic properties can be observed or measured e.g. state, colour, temperature, pressure) * Concentrations of reactants and products are constant but not necessarily equial * Continual microscopic change between reactants and products * Rate of forward reaction equal to backward reaction occurring simultaneously (equilibrium is dynamic) * Chemical reactions can either be reversible or irreversible * Irreversible reactions: when all the reactants turn into products and goes into completion. Indicated by a one way arrow i.e. complete conversion of reactants to products when stoichiometric amounts of reactants are mixed * Reversible: when reactants can form products (forward reaction) and products can form reactants (reverse reaction) indicated by a two way arrow. It will reach an equilibrium
What is the purpose of the conc acid in esterification
* Concentrated sulphuric acid acts as a catalyst by proving an alternate pathway which reduces the activation energy (Ea) required hence increasing the rate of reaction Increase the yield -- * Concentrated H2SO4 acts as a dehydrating agent as it has a high affinity for water * This causes the water to be removed from the reaction and th equilibrium shifts to the right in order to increase the concentration of water (LCP). This increases the yield of ester
End point vs Equivalence point
* Equivalence point — occurs when stoichiometric ratios of acids and bases have neutralised each other, leaving only an aqueous alt in solution * Salt can be acidic, basic or netural depending on the acids and bases used * At equivalence opint the pH changes very rapidly as a drop of acid/base is enough to significantly change the concentration fo hydrogen ions. Hence the titration curves are very steep at equivalence point * Titration with a weak acid and a weak base is conducted with a pH meter instead of an indicator as the colour change occurs over a large range making it difficult to determine end point * When the indicator permanently changes colour * In an ideal titration, the end point is the equivalence point and the difference is the titration error. Hence an appropriate indicator must be selected such that its end point, when it changes colour, is at its equivalence point * Acidic salt solution = methyl orange (3.1-4.4) * Neutral salt solution = bromothymol blue (6.0-7.6) * Basic salt solutions = phenolpthalein is used as it changes pH around 7 (8.3-10.0)
Describe the uses of esters
* esters occur naturally in living thigns * they ahve strong flavours and odours so they are used for fragrances and falvours in processed foods * they are cheaper than natural extracts and pose little health hazards USES * flavourings: pentyl ethanoate as pear flavourings * cosmetics i.e. perfumes i.e. phenyl ethanoate as jasmine fragrance * industrial plasticiser i.e. dialkyl phthalate as a PVC plasticiser *
What is the effect of a catalyst, noble gas on equilibrium
* he addition of inert gases increases the pressure of the system, but it does not affect the partial pressure, hence there is no change in the position of equilibrium. The addition of inert gases reduces the amount of successful collisions between other gas moleucesl, decreasing the rate of both the forward and reverse reaction * If the no of mols of gas sis equal to the no of mols of products, then change of pressure will not affect position of equilbrium * A chemical substance which provides an alternative pathway for reaction by decreasing the rate of reaction by lowering activation eenergy * This will increase he rate of both the forward and reverse reaction without affecting the position of the equilibrium
Why is refluxing needed in esterification and what is it? What are other features of esterification
* refluxing is the heating of the reaction mixture and the cooling of the volatile reactatn gases using a reflux condenser during esterification * the alternative to refluxing is using a closed container however this has the risk of building up pressure --> explosion therefore not safe WHY IT IS NEEDED 1. allows the reaction to be carried out at a higher temperature. This is becasue esterification is a slow reaction tehrefore higher heat is needed to increase the rate of reaction. But since the reactants are volatile, they can escape the vessel therefore the cooling of these reactants ensures no escape 2. to ensure no volatile gases escape. The condensor avoids the loss of any material from the reaction vessel. Cold water circulates the condesor to cool the vapour, condensing the liquid and making it fall back into the reaction mixture 3. For safety because the escaping volatile gases can be flammable FEATURES OF CONDENSOR * The reflux condenser is a normal condenser placed vertically inside a round bottom flask and attached to a retort stand by a boss head and clamp * The top of the reflux condenser is open to ensure no pressure builds up inside * The water in is below the water out to ensure the condenser fills completely instead of run down if it were at the top OTHER FEATURES * boiling chips which encourage the mixing of the reactants to evenly distribute the heat. They prevent 'bumping' by providing a large SA for vaporisation which would prevent the superheating of the liquids and explosive ejections of the vapours I.e. rigorous boiling or bubbling * the water bath allows for even heating and greater safety as organic compoudns are voltage and flammable which means must be kept away form naked flames. ALLOWS FOR NOT USING A BUNSEN BURNER WITH THE HOT PLATE
Assess the evidence which indicates increases in atmospheric concentrations of oxides of sulfur and nitrogen
* since the industrial revolution, there has been an increase in the oxides of sulphur and nitrogen due to th combustion of coal, fuels in cars and in extracting metal * the first sign of increase in these levels was in the 18th entury when it was noted taht there were marine life unable to be supported in fresh water environements * in the 1950s and 60s, air quality began to significantly decrease due to pollution episodes in London and USA, which cuased many deahts. This led to control emissions form factories to improve the air quality * Acid rain is a bigger problem in Europe, China and Russia as they combust sulfur containing coal for electricity. SULFUR DIOXIDE AND EVIDENCE * comes from the combustion of sulfur containing coal for electricity and the smelting of metals ZnS + O2 --> SO2 + ZnO * the evidence comes the analysis fo ice cores samples showing increased concentration fo sulphates and also the greater acidificaiton of lakes in North America and europe * also there has been an increase in the destructive effects of acid rain on limestone strcutres e.g. the Taj Mahal CaCO3 + H2SO4 --> CaSO4 + CO2 + H2O * currently, the levels have declined in western europe and north america however it still remains an issue in Asia, Eastern Europe and South america where industrial coal use is prevalent NITRIGEN OXIDES AND EVIDENCE * comes from the interanl combustion of fuels in cars which release nitrogen and the coal burnign power stations N2 + O2 --> 2NO * int he presence of sunlight it is oxidised into NO2 NO + O(radical) --> NO2 * the evidence comes from the increased photochemical smog within industrial areas/cities and incresae of acid rain * also there have been quantitative analysis of the gases in the atmosphere which shows increased levels of NOx * most cities have not improved or deterioriated because the increase of vehical kilometres and city populations have cancelled out the benefots of incresaing emission control * to decrease the motor cars have emission controls in order to decrease the emission of nitrogen Ice core samples in Antarctica have 10 percent increase of N2O SIGNS OF IMPROVEMENT * Due to the US EPAs Acid Rain Program, the US has a 33 percent from 1982 to 2002 decrease in the emissions of sulfur due to the flue gas desulfurisation * New fuel additive catalyst e.g. ferox is used in gasoline and diesal engiens to lower the emisiosns fo SO2 gas in the atmosphere Difficulties in obtaining evidence * They exist in small concentrations of 0.01ppm * Infrared spectroscopy used to measure the concentration not available until 1970s — only had a few decades in which we can measure the concentrations * Sulfur oxides and nitrogen oxides may exist in aqueous forms in hydrosphere and biosphere which is hard to measure EVALUATION * Assessment: Overall there is sufficient evidence to suggest an icnrease in atmospheric concentration of SO2 and nitrogen oxides since the Industrial revolution despite the difficulties involved in obtaining evidence * The quantitative measurements (analysis fo ice core samples and direct measurements of gas concentration in the atmosphere) are accurate and reliably point to increases in concentration of NOx however the technology for measuring low concentration has only been available in recent times * Qualitative evidence for smog and acid rain is less reliable, as these issues could be caused by other acidic oxides and other factors may have led to their increase
Describe how the self ionisation of water contributes to the pH scale and how to determine degree of ionoisiation
* the water undergoes self ionisation which results in a very small concentration of H3O+ and OH- in the water, however since the ratio is 1:1 they will neutralise each other out to form neutral water 2H2O <-> H3O+ +OH- * the equilibrium constant is given by the product of the hydronium and hydroxide ions. Since it is an endothermic reaction then the higher the temperature, the higher the value than if it was cooler At 25 degrees: Kw = [H3O+][OH-] = 1 * 10^-14 pH = 14 + log[OH-] Since pOH = -log[OH-] Then pH + pOH = 14 Strong acids 100 percent ionisatin except weak no: [H3o+]/[HA] * 100
Compare acetic, citric, hydrochloric and sulphuric acid (including pH and degree of ionsiation)
0.1M SOLUTOIN: Acetic * ethanoic acidd CH3COOH * monoprotic * 1.3 percent ionisation and pH = 2.9 * use: vinegar CH3COOH <-> CH3COO- + H+ Citric * triproctic * 8.7, 0.13 and then lower percentage of ionisation and pH = 2.1 * weak * fruits * C6H8O7 * 1-hydroxypropane- 1,2,3 - tricarboxylic acid (3 equations in water) C6H8O7 <-> C6H7O7- + H+ Hydrochloric acid * * Chemical formula: HCl * Monoprotic acid * Strong acid * Stomach acid HCl —> H+ + Cl- 100 percent and pH = 1
Preparation of a natural indicator
Aim: To prepare and test a natural indicator (red cabbage) Hypothesis: A natural indicator will be prepared by extracting he pigment (anthocyanin) from red cabbage which will change colours when tested with acidic, natural and basic substances 1. Place shredded cabbage in 250mL beaker and just cover with distilled water (about 50mL). Slowly boil leaves until water turns dark reddish purple and the leaves lost most their colour (anthocyanin pigment) 2. Allow to cool and pour liquid into clean 250mL beaker 1. 2 drops of 0.1M NaOH, 9.1M HCl and of each of the pH solutions onto separate spots on the spot testing plate 2. Add a few drops of anthocyanin pigment from cabbage to each sample until definite colour change is observed 3. Record the colour of indicator 4. Repeat Results * In a neutral solution, the red purple cabbage indicator remained purple * In an acidic solution, the red purple cabbage indicator turned red/pink * In a basic solution, the red purple cabbage indicator turned blue/green/yellow From strongly acidic - neutral - strongly basic: Red —> pink —> purple —> blue —> green —> yellow Justifications * Red cabbage was used as the vividly coloured red purple pigment (anthocyanin) can be easily extracted and used as a natural indicator Limitations * The natural indicator cannot determine the exact pH as colour change is subjective * The natural indicator (anthocyanin) is a weak, diprotic acid which affects the pH of the solution * A pH meter is sued to give a more accurate reading as well as not affecting the pH of solution Safety * Safety glasses and clothes are used to protect the eyes and body from contact with acidic or alkaline substances which may cause blindness Validity * Variables controlled: same red cabbage, same observer judging the colour change and same equipment used * Independent variable: common household substances * Dependent variable: colour change of indicator Reliability * The reliability of the experimental procedure could be improved by obtaining similar results through repetition Accuracy * The experimental procedure was accurate as it changed colours in acidic, neutral and basic solutions, which is close to the expected colour * The accuracy could be improved by using more precise measuring instruments such as a pH meter Conclusion: A natural indicator (red cabbage) was successfully prepared and tested with a range of acidic, neutral and basic substances which resulted in a colour change
Natural indicator practical
Aim: to prepare and test a natural indicator Hypothesis -- a natural indicator will be prepared from the anthocyanin pigment within the red cabbage, which will change colours upon tested with a range of substances with different pH Method - chop up the red cabbage and place in water - heat with a hot plate until the red purple pigment (anthocyanin) is extracted. Then filter off --> leaves behind natural red cabbage indictor - place the same amount into different test tubes and then add a range of substances observe tis colour changes e.g. CH3COOH and H2SO4 Observations Acidic -- red/pink Neutral -- purple Baisc -- blue, green, yellow Justifiations -- used red cabbage because the pigment is easily extracted and can change colour when added to a range of substances Limitation — Cannot determine exact pH ( colour change subjective) — Anthocyanin —> a weak diprotic (lose 2 hydrogens) acid —> affect pH of solution
Identify amphiprotic substances and construct equations to describe their behaviour in acidic and basic solutions
Amphiprotic -- can act as both a proton donor and acceptor. All amphiprotic substances are amphoteric, but not all amphoteric are amphiprotic (these are lewis acids and bases which accepts and donates electron pairs respectively) E.g. with water: HCO3-(aq) + H3O+(aq) <-> H2CO3 + H2O Basic: HCO3- + OH- <-> CO3^2- + H2O
Describe the process of titration and the cleaning of the equipment
Before hand, ensure to wash the pipette, burette and conical flask with distilled water from a wash bottle 2 times to avoid contamination. Wash the pipette and burette with the solution going inside it to make sure that the concentration is preserved and no dilution occurs 1. Set up burette attached to retort standard by burette clamp as shown aboave 2. Fill burette with a solution (usually titrant) until bottom of meniscus is on zero mark 3. Draw up exactly 25mL of solution (this is known as an aliquot, usually this solution is the analyte) using a Pasteur pipette and transfer into conical flask by resting tip against side of glass to let last drop run down * Do not blow the last drop off solution. The pipette already accounts for this volume * To use, deflate the bulb and then press 2 to fill up, and then empty liquid by pressing____ 4. Add 3 drops of suitable indicator into conical flask and swirl the flask to mix it into the solution * Only a few drops added because indicators are weak acid/base which may affect pH of solution in conical flask 5. Place conical flaskk on white tile or background used to observe colour change, and under the tap of the burette 6. Stream gently the solution from the burette, swirling the flask. When the colour starts to disappear more slowly, add solution in drop wise. When colour change permanent, stop and take measurement of volume by reading the bottom of the meniscus. 7. This is the rough titre. Don't include 9. Repeat steps 2 to 8 theree times including re rinsing and determining the average titre: volume of titrant used to reach endpoint
Describe the use of the pH scale in comparing acids and base
The pH scale is a logarithmic quantiative scale of base 1- which emasures the concentartion fo hdyrgoen ions to determine the strength sof acids and bases in aqueous solution. Therefore one change in pH = 10 fold change in concentration pH = -log[H+] [H+] = 10^-ph
Plan and perform first hand investigation to measure the pH of identical concentrations of strong and weak acids
DRAW A DIAGRAM 0.1M of acetic acid vs HCl Beakers and pH meter Place 100mL of each solution into different beakers Remove the pH meter from the buffer solution and calibrate, rinse with distilled water and then place into acetic. rinse with distilled water, restandardise before putting into HCl REUSLTS HCl was 1 and CH3COOH was 2.9 JUSTIFICATIONS * since pH is the measure of the ocncentrations of the hydrogen ions, then by conconctrolling the concentration fo the acid, the strength can be determined. ** HCl had a higher degree of ionisation releasing more hydrogen ions as it is a strong acid with its ionisation going to complete reaction HCl(l) —> H+(aq) + Cl-(aq) where ΔH < 0 * Acetic acid has a low degree of ionisation releasing fewer hydrogen ions as it is a weak acid with its ionisation exaction in equilibrium. This is because weak acids ionise less readily in water and require a much higher concentration to equal that of a stronger acid * Note: while there are degrees of weakness of acids, there are no degree of strength for strong acids. If an acid is strong it completely ionises but if it is weak then it partially ionises Limitations — none Safety * HCl is a strong acid which is corrosive so protective clothing and safety glasses must b warn Validity * pH is a measure of acidity and depends upon strength and concentration * Therefore the experimental procedure was valid as variables were controlled such as: same concnetrationsof acid, same pH meter used and rinsed between tests and same temperatrue * Independent: weak acid and strong acid, CH3COOH and HCl respectively * Dependent: pH of each aicd Reliability * Reliable as similar results were obtained upon repetition Accuracy * The experimental procedure was accurate as the pHs recorded were the expected value and the pH of a stronger acid was lower than the pH of the weaker acid Conclusion: The pHs of strong acids (HCl) are lower then the pH of weak acids (CH3COOH) with identical concentrations because they have a higher degree of ionisation
Qualitatively describe the effect of buffers with reference to a specific example in a natural system
Definition: A buffer is a solution of a weak acid and its conjugate base (e.g. CH3COOH and CH3COO-) or a weak base and its conjugate acid (e.g. NH3 and NH4+) * They serve to resist the rapid changes in pH when a base or acid is added * In a normal, natural aqueous solution, the addition fo a small amount of 0.1M strong acid will change the pH by factors of hundreds of thousands. With a buffer, the addition of the same amount of acid would not change the pH appreciably * Buffers are important in living, natural systems as they can maintain pH within a certain range via LCP HA + H2O ⇌ H3O+ + A- Carbonic acid/hydrogen carbonate buffer in blood * Blood consists of a carbonic acid buffer system to maintain a pH of 7.4. This pH is necessary for enzymes to function optimally, cell function and metabolism H2CO3(aq) + H2O(l) ⇌ H3O+(aq) + HCO3-(aq) ΔH < 0 Explanation that is not using LCP * When an acid is added, a neutralisation reaction will occur between the buffer and the acid H3O+(aq) + HCO3-(aq ⇌ H2CO3(aq) + H2O(l) * When a base is added: OH-(aq) + H2O(l) ⇌ H3O+(aq) + HCO3-(aq) Explanation using LCP * When an acid is added (acidosis) e.g. increase expiration of CO2 this is a disturbance to the equilibrium thus the it will shift to left in order to minimise the disturbance, increase the reaction of the reverse reaction and therefore forming water and carbonic acid (LCP). This is to decrease H3O+ H3O+(aq) + HCO3-(aq ⇌ H2CO3(aq) + H2O(l) * When a base is added e.g. OH- such a decreased expiration of CO2 (alkalosis), the position of the equilibrium will shift to the right to decrease the OH- OH-(aq) + H2O(l) ⇌ H3O+(aq) + HCO3-(aq) * Hence the hydrogen carbonate iron adn the carbonic acid act as buffer to resit pH changes and maintain a pH of 7.4 * However when when too much acid or base is added, there will not be enough hydrogen carbonate ion or carbonic acid to neutralise the change * The maximum amount of acid or base the buffer can neutralise before changing pH is called the buffer capacity
What is esterification
Esterification is a condensation reaction between an alkanol and alkanoic/carboxylic acid, forming an alkyl alkanoate and water. It is a slow endothermic reaction and is reversible, usually occurring in the presnecne of a ocnc acid. At room temp the equilibrium lies to the left * the OH in the water is from the carboxylic/alkanoic acid * the other H from the alkanol * Note: the oxygen atom which forms the bridge btween the alkanol and alkanoic acid is from the alkanol,
Draw a diagram of strong acid vs weak acid
First diagram with all the molecules not ionised yet
Write ionic equtions to represent the ionisation fo acids
Hydrochloric acid HCl(l) + H2O(l) —> H3O+(aq) + Cl-(aq) Nitric acid HNO3(l) + H2O(l) —> H3O+(aq) + NO3-(aq) Sulfuric acid H2SO4(aq) + H2O(l) —> H3O+(aq) + HSO4-(aq) HSO4-(aq) + H2O(l) ⇌ H3O+(aq) + SO42-(aq) Acetic acid (ethanoic acid) CH3COOH(s) + H2O(l) ⇌ CH3COO-(aq) + H3O+(aq) Citric acid C6H8O7(s) + H2O(l) ⇌ H3O+(aq) C6H7O7(aq)-
What are the substacnes needed for esterification and write out a word equation. Give example
alkanol + alkanoic acid <-> alkyl alkanoate + water methanol + pentanoic acid <-conc H2SO4-> methyl pentanoate + water
Computer based technologies
Measure vinegar and HCl Hypothesis: The concentration fo vinegar will be determined by titrating it against a NaOH secondary standard and using a pH meter to determine equivalence opint Equipment * Vinegar * Standardised NaOH 0.1M * 0.1M HCl * pH probe/data logger * Computer * Magnetic tirrer and stirring bar * Distilled water * Titration equipment Method A: Titration of SA and SB 1. Transfer 25mL aliquot of 0.1M HCl into a clean, dry 100mL beaker. 2. Fill the burette with 0.1M NaOH 3. Calibrate and rinse the pH meter before immersing into analyte solution 4. Run 1mL of NaOH into the beaker. Measure pH 5. Repeat step 3 until the pH starts to change rapidly 6. Add 0.5 mL at a time measuring the pH after each addition. 7. When the pH change becomes less, then continue adding NaOH in 1 mL aliquots until the pH graph becomes almost horizontal (equivalence point. This is indicated by a large pH change with a small amount of titrant). 8. Repeat steps 1-7 Method B: Titration using a strong base and a domestic acid (vinegar) 1. Set up equipment as above 2. Using a clean, dry pipette transfer 10mL of vinegar into clean dry 100mL volumetric flask 3. Fill flask to mark with distilled water, stopper the flask and mix the solution by inviting the flask a number of times. This is a 10 percent vinegar solution 4. Use a clean pipette an transfer 25mL aliquot into 100mL beaker 5. Calibrate and rinse the pH metre before immersing it in the analyte solution 6. Titrate the vinegar with 0.1M NaOH standardised solution until equivalence point is reached (indicated by a large change in pH with small amount fo titrant). It should be above 7 as a strong base and a weak acid is being titrated 7. Record the tire 8. Repeat steps 2-7
Name the 4 indicators
Methyl orange 3.1 to 4.4 Red to Yellow Bromothymol blue 6-7.6 yellow to blue Phenolpthalein 8.3 to 10.0 Colourless to pink Litmus red to blue 5 to 8
Identify data, father and process information form secondary sources to identify examples of naturally occurring acids and bases and their chemical composition
NATURALLY OCCURING ACIDS * citric acid (formula) found naturally in fruits * acetic acid is found naturally in the oxidation fo ethanol which is derived from the abcterial fermentaiotn fo sugars * hydrochloric acid is found in the stomach and it helps digestion NATURALLY OCCURING BASES * metallic oxides are found naturally in rocks * ammonia (NH3) is found naturally in urine * calcium carbonate is found naturallly in limestone
Analyse information form secondary sources to assess the use of neutralisation reactions as a safety measure or to minimise damage in accidents or chemical spills
Neutralisation reactions are used to minismie the dmaage in accidents or chemical spills as acids are corrosive and bases are caustic The chemicals msut be: * amphiprotic * weak * powdered (solid) * easily transportable * cheap Since neutarlisation reactions are exothermic, if it occurs on the skin, it msut be washed off with water immedietly as an exothermic reaciton may cause burns FOR EXAMPLE Anhydrous NaHCO3 Acid spills: NaHCO3 + H+ --> Na+ + H2O + CO2 Basic spills: NaHCO3 + OH- --> Na+ + H2O + CO3 2- Na2CO3 can be good for acidic spills. Even in excess it is safer than NaOH or lime in excess Na2CO3 + HCl --> H2O + CO2 + NaCl If large spills occur, sand is used to prevent acid from running off into drains or soils on the side of the road. ONce all the acid is absorbed, the sand is sent off to be neutralised and Na2CO3 pwoder + water neutralise and dilute any more acid
Describe an experiment you did to prepare an ester. Risk, isolation, justification. DRAW DIAGRAM
Reaction between 1-pentanol and acetic acid Equipment -- hot plate, concentrated H2SO4, conical flask, stopper, separating funnel, boiling chips Chemicals -- 10mL 1-pentanol, 12mL glacial acetic acid, 1mL of 9M concentrated H2SO4, 15mL of 1molL^-1 Na2CO3, anhydrous calcium chloride 1. add the acid, pentanol, H2SO4 into round bottom 150mL flask 2. Add boiling chips and assemble like diagram (book), don't use bunsen burner as this is very dangeorus 3. Connect tubing to tap and condenser and turn on the water 4. Heat mixture on hotplate for 30 minutes and turn off the water 5. Remove the flask and pour contents into a separating funnel with 10-15mL of water. Stopper the funnela nd shake, allowing layers to separate, drain off and discard the lower aqueous layer 6. add Na2CO3, shake, drain and discard loewr layer so ester in separating funnel 7. Smell ester and record the scent. Pour mixture into flask and add CaCl2 and distill the ester Justifications Isolation of the ester * can isolate the ester via disitllation by assembling a distillation apparatus and placing the solution into a distilling flask, boiling chips * distil the ester * the receiver msut be cooled in an ice bath. As esters have the lowest BP int he vessel, it will come out first (must have the highly volatile substance stay inside) * to further purify, add Na2CO3 prior to distllation because this will allow the oily organic layer to be separated using a separating funnel (above) * distilling this will further purify the ester RESULTS (name the products) JUSTICATIONS * the reflux condeser used to allow for use of high termpeatures for increasing teh rate of reaction, preventing the loss of any voltaile reactants + safety reasons * boiling chips allow for the reactatns to mix and gentle heating, and prevents rigorous boiling or bubbling * the use of H2SO4 (etc etc) SAFETY * the conc H2SO4 is highly corrosive thus when not in use, must close the container. Safety glasses too * waft the smell don't directly place nose next to the reaction vessel as it can cause respiratory issues * hot plate
Identify and describe some everyday uses of indicators including the testing of soil acidity/bascity
SOILING TESTING pH of soil needs to be adjusted to ensure optimal plant growth Sample is collected and barium sulphate which is white is added this is so easier to observe the colour change The indicators are added to determine the asicidity or basicity and compared to a pH chart (univrsal indicator) or using the pH Too acidic -- add CaCO3 (lime) Too basic -- add ammonium chloride to lower the pH * NH4Cl * NH4+ + H2O ⇌ NH3 + H3O+ SWIMMING POOLS * The pH of swimming pools needs to be maintained at about 7.4 for sanitation, to avoid irritation of the skin and to avoid the corrosion of metals 1. Sodium hypochlorite is added to kill the microbes. OHCl is the active form which kills microbes but is relatively unstabe * The presence of OH- ions makes the water basic, so HCl is added to return neutrality. This must be done to protect the eyes and throats of swimmers from becoming irritated 2. Sample of water is collected 3. Phenol red with a pH range of 6.8 to 8.4 is sued as an indicator to determine the acidity or basicity * Chemicals may then be added to adjust the pH of the swimming pool
Compare probes vs indicators and how to standardise the pH meter
STANDARDISING * ensure the probe to be in distilled water 24 hrs before use at least. Take it out, dry and place into a buffer solution (usually arond 4, 7 or 10) * before measuring: remove from teh buffer solution, rinse the electrodes with distilled water, place the electrodes into the unknown and read pH * must alwyas do this when changing between solutions * Indicators * Are weak acid/base and are a destructive method of testing becasue they can change the pH * Only give a pH range therefore less accurate * Cheaper than pH emters * pH meters * Do not alter pH and a non destructive method of testing * Given a value for pH reading and are more accurate * More expensive than indicators
What are natural and industrial sources of sulfur dioxide and nitrogen? What are the dangers and how is it minismeidd/
SULFUR Natural * volcanoes release sulfur S + O2 --> SO2 * geothermal hot springs release directly * the decomposition of organic matter with bacteria H2S + O2 --> SO2 + H2O Indsutrial * sulfur rich metal ores when extracting metals * 2ZnS(s) + 3O2(g) —> 2ZnO(s) + 2SO2(g) * the combustion of fossil fuels of sulphur rich coal S(s) + O2(g) —> SO2(g) * from the contact process NOx Natural * lightning produces high temperatures in the atmosphere, cuase the N2 to react with O2 and for NO, which can oxidiseinto NO2 * nitrogen fixing bacteria in soil N2 + O2 --> N2O * oxidation of NO in sunlihgt forms NO2 Unnatural * power stations produce high tempeatures (reactions below) * * Internal combustion car engines produce high temperatures which combines to atmospheric nitrogen and oxygen to form NO, further oxidising to NO2 * N2(g) + O2(g) —> 2NO(g) * 2NO(g) + O2(g) —> 2NO2(g) CONCERNS * photochemical smog can form from NOx which is a respiratory irritant. Bad for asthamtics (from NOX and hydrocarbosn combining with oxygen in the presence of sunlight forming ozone) * the NO2 in the presence of UV will form NO and oxygen radical. This radical can form ozone gas wwhen reacted to oxygen gas * acid rain REDUCTION * catalytic convertors in the car engines makes the NO convert back to N2 NO + CO --> N2 + CO2 * the SO2 can be removed through scrubbers which pass the gaseous emissions through Mg(OH)2. The magnesia sulphite that is produced is disposed in landfill or conerted back to Mg(OH)2 * SO2 produced in the smelting of metal sulphides can be collected and used to make H2SO4
What is pH determined by
Strength and concentration
Compare strong vs weak acids, concentrated vs dilute acids
Strong = 100 percent ionisation of H+ ions to form hydronium ions Weak = partial ioniation of the + ions Conc = more molecules per unit of solutoin (independent of ionisation) Dilute = low number of acid moleucels per unit fo solution
Describe the relationship between an acid and its conjugate base and a base and its conjugate acid and relative strengths. Provide examples
Strong forms neutral/very very weak. This is becasue it has a small tendency to accept/donate the proton Weak forms weak -- moderate tendency to accept/donate proton Neutral -- forms strong. Large tendency to accept proton E.g. HCl forms a neutral Cl- CH3COOH forms CH3COO- weak base OH- is a strong base which forms neutral acid of water Neutral acid forms strong base of OH-
Describe the differences in MP and BP for alkanes, alkanols, esters and alkanoic acids
The differences in MP and BP depend on the relative strenght of the intermolecular forces Alkanes --> esters --> alkanols --> alkanoic acids ALKANES * alkanes contain weak dispersion forces as they are non polar, therefore they have a low BP and MP. Formed from the movement of the electrons which form tempoerary dipoles * molecular weight increases the BP increases. Note that the MP will not not increase with this trend, as it also depends on the shape of the moleculse and how ti forms the H bond in a reggular arrangement in a solid. ESTERS * due to the presence of the C-O and C=O group it is polar, as it forms dipole dipole interactions between the molecules. This means it will have a higher BP than the alkanes * as molecular weight increases, the dispersion forces increases therefore teh higher BP ALKANOLS * due to the presence of the O-H and C-O bond it can form H bonds and dipole dipole forces between molecules (H bond = hence the higher solubility of the shorter chains) * due to stronger interactions with other moleucles = higher BP * stronger dispersion forces also because of the presence of the oxygen atom * as the chain length incrases, the solubility decreases * higher molecular weight = more dispersion forces = higher BP ALKANOIC ACIDS * highest because of the presnce o teh C=O, C-O and O-H bonds * higher than alkanol as it forms more extesnion hydrogen bonding and dipole dipole interactions due to the presnece of the 2 oxygen atoms, forming stable dimers * they exist as dimers iwth 2 strong hydrogen bonds between each molecule * stronger dispersion forces when increase molecular weight = higher BP
WHAT IS TITRATION Describe the process of making a standard solution and what properties are wanted with a primary standard (examples included)
Titration -- using a neutralisation reaction to reach end point in order to determine the concentration of a substance with known volume using another solution with known concentration and volume -> EXOTHERMIC A primary standard is a pure, stable substance used to make a standard solution Solid -- prevents evaporation and also easier to weigh Non hygroscopic and non efflorescent -- does not absorb water and not release water respectively High molecular weight -- reduces error when weighing High solubility -- ensures all dissolved High purity -- no chemical impurities which would cause an inaccurate mass Chemically stable -- doesnt react with oxygen in the air or CO2 which owuld produce an inaccuraate mass Suitable -- anhydrous Na2CO3, oxalic acid crystals, NaHCO3 Unsuitable -- NaOH (hygroscopic and reacts with CO2 in the air), HCl (anhygroscopic) and H2SO4 (hygroscopic) To prepare a 250mL standard solution of 0.05M 1. Pick a suitable primary standard * If analyte (substance your analysing) is base, use suitable acidic primary standard e.g. oxalic crysclese (COOH)2*2H2O * If it is an acid, use suitable basic primary standard e.g. anhydrous sodium carbonate (Na2CO3) 2. Wash a 250mL volumetric flask, small beaker, stirring rod, glass funnel and Pasteur pipette using distilled water from a wash bottle * This ensures no contamination and standard solution is accurate 3. Calculate mass of primary standard needed to make 250mL of 0.05M standard solution using C=n/V 4. Carefully weigh the required mass in a small beaker using an electronic balance 5. Place 200mL of distilled water from the wash bottle and use stirring rod to ensure that all solute has dissolved 6. Transfer primary standard from small beaker into volumetric flask via funnel by washing it with distilled water, ensuring all solute in beaker and funnel is washed into the volumetric flask (for accuracy) 7. Remove the funnel and pipette (Pasteur pipette) the rest of the water in until the bottom of the meniscus is on the 250mL mark (if you add too much, start again) * You pipette the water because the neck of the flask is very skinny, thus making it more accurate * Removing the funnel will slightly increase the water level as water trapped in the neck of the flask will run down into solution * You prepare it in a beaker so you can stir it * If you add slightly too much, start the process again 3. Add a stopper and invert the flask.3 times 4. Label the volumetric flask with the name of solution, concentration and date and your name
Decarbonating soft drink prac
To decarbonate a soft drink and measure the mass changes involved and calculate the volume released at 100kPa at 25 degrees celsius Hypothesis: Opening the soft drink will decarbonate it, and the mass lost by the soft drink willequalt he mass of escaped carbon dioxide, which then can be used to calculate the volume of gas 2 METHODS 1. Letting the soda go flat * weigh soda bottle before
Solve problems and perform a first hadn't investigation to use pH meters/probe sand indicators to distinguish between acidic, basic and neutral substacnes Choose equipment and perform a first hand investigation to identify the pH of a range of salt solutions
To use pH meters and indicators to distinguish between acidic, basic and neutral substances * the pH meter will be more accurate Lemon juice, distilled water, bicarb soda, acetic acid/white vinegar, washing powder, dish washing liquid, universal indicator, soil * need pH meter, indicators, testu tube rack 1. Remove pH meter from buyer solution and rinse electrode distilled water form wash bottle. Turn the meter on and place electrode in same, take reading rinse the pH electrode. Each time the pH meter is being used, rinse wiht distilled water 2. Take small sample of 10 different acidic, basic and neutral substances 3. Place small amount of solution to be tested in small beaker. Insert pH meter and record 4. Measure the pH by using a variety of indicators and use a colour chart to narrow pH 5. Record results Results * At 25 degrees celsius the acidic substances had a pH<7, neutral pH=7 and basic pH>7 * There was a small difference between pH measured by the pH meter and the pH determined by the indicators Safety * Acidic substances are corrosive and basic substances are caustic so protective clothing and safety glasses must be worn * Fragile glass of test tube — let teacher handle it, keep glass away form the edge of the tbale Validity * Variables controlled: same amount of substances tested, same measuring equipment used, pH meter standardised, pH meter rinsed using distilled water after each substance was tested to avoid contaminiation * Independent: the substance tested * Dependent: the Ph of substance Reliability * Reliable as similar pH readings were obtained when measured with indicators and a pH meter Accuracy * Te experimental procedure was accurate as the pH mdetermiend of each substance was close to the real pH * More precise measuring instruments such as pH meter instead of indicaots improved accuracy Conclusion * pH meters/probes and indicators were successfully used to distinguish between acidic (pH<7), basic (pH>7) and neutral (pH=7)
Cleaning equipment for titration
Volumetric flask and stopper Function: Prepare and hold SS of known concentration Rinsing: distilled water to avoid contamination and preserve moles of primary standard. It is left wet and stopper used to prevent evaporation Conical flask Function: hold reactants during titration, usually analyte. Its shape prevents reactants from spilling as it is swirled Rinsing: dialled water to avoid contamination and diluting it slightly preserves number of mols in primary standard in solution Pasteur pipette Function: to measure and transfer exact volume (usually 25mL) os solution (usually analyte) into conical flask. A pipette bulb/filler used to draw up liquid * To use, press 2 to fill the pipette and then let the bottom of the meniscus be at the calibration line Rinsing: after rinsing with distilled water 2 times, Piette is rinsed with solution 2 times going to measure to ensure there is no dilution and preserve concentration of solution * To rinse attach pipette filler to pipette to suck up some of the solution to a level just not the bulb. Remove the pipette filler and then rotate the pipette as the solution is allowed to run into the sink Burette Function: to measure and dispense an accurate volume of solution (titrant) into the conical flask until end point is reached * To use, Rinsing: after rinsing with distilled water 2 times, burette is rinsed with solution ti is going to measure 2 times to make there is no dilution and to preserve concentration of titrant * To wash, add 5-10mL of solution via the top of teh burette using a small funnel. Take the burette out of hte burette clamp and let solution run into sink while rotating the burette
Explain the formation and effects of acid rain
pure water has pH of 7 whilst rain water has pH of 6 to 6.5 due to CO2 in the atmosphre dissolving with teh water to form carbonic acid CO2 + H2O --> H2CO3 H2CO3 + H2O --> H3O+ + HCO3- Acid rain below 5 SULPHURIC ACID RAIN S + O2 --> SO2 SO2 + O2 --> SO3 SO3 + H2O --> H2SO4 __ SO2 + H2O --> H2SO3 H2SO3 + O2 --> H2SO4 NITRIC ACID N2(g) + O2(g) à 2NO(g) 2NO(g) + O2(g) à 2NO2(g 2NO2(g) + H2O(l) à HNO2(aq) + HNO3(aq) * nitrous acid further oxidises into H2SO4 EFFECTS * lowers the pH of small bodies of water affecting marine life > the lowered pH can lead to the leeching of phosphates and sulphate ions from fertilisers, which cause algal blooms. This hinders the supply of oxygen for the marine life > can cause the leeching of heavy metals from the soil into the lakes killing aquatic life > can dissolve the calcium in the shell fish and also kills eggs Damages to the soils and plants * can cause the defoliation of the leaf by corroding it, removing its cuticle. This makes it more prone to infetions and insect * it can lower the pH of the soil which means it will dissolve some of the minerals affecting plant growth. This also makes it ahrder to absorb water. Trees in mountainous regions have to deal with the acidic fog and clouds Buildings * Damages to buildings and status * Acid rain cords marble, or limestone, buildings and statues (anything with calcium carbonate) CaCO3(s) + H2SO4(aq) à CaSO4(aq) + CO2(g) + H2O(l) * It also corrodes metal buildings and statues Fe(s) + H2SO4(aq) à FeSO4(aq) + H2(g)
